ATOMIC STRUCTURE

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Atomic Structure

SCIENCE AND TECHNOLOGY

Notes

MODULE - 2Matter in our Surroundings

5

ATOMIC STRUCTURE

In lesson 3, you have studied about atoms and molecules as the constituents of matter.You have also learnt that the atoms are the smallest constituents of matter. In lesson 4you studied about the chemical reactions, their types and the ways to represent them.You know that according to Dalton’s atomic theory, the atoms of different elementsare different and in chemical reactions the atoms are rearranged between differentreacting substances. However, today we know that the atom is not indivisible as itwas thought by Dalton. The atom has a structure and contains smaller constituentsin it. In this unit, we would attempt to find out the answers to some of the questionslike, “What is the structure of an atom?”, “What are the constituents of atoms?”,“Why the atoms of different elements are different?” and so on.

We will begin this unit with the study of the discoveries of sub-atomic particles suchas electron, proton etc. Then, we will take up various atomic models proposed onthe basis of these discoveries. We will discuss how various models for the structureof atom were developed and also explain the success as well as the shortcomingsof these models. This will be followed by the description of the arrangement or thedistribution of electrons in the atom. This arrangement is known as electronicconfiguration. These electronic configurations are useful in explaining variousproperties of the elements. These also determine the nature of chemical bonds formedby it. This aspect is dealt with in lesson 7 on chemical bonding.

OBJECTIVES

After completing this lesson, you will be able to:

� recall the evidences showing the presence of charged particles in matter;

� describe the discovery of electron and proton;

� explain Dalton’s atomic theory and its failure;

� discuss Thomson’s and Rutherford’s models of atom and explain theirlimitations;

Atomic Structure

SCIENCE AND TECHNOLOGY 94

Notes


� explain the Bohr’s model of atom (in brief);

� describe the discovery of neutron;

� compare the characteristic properties of proton, electron and neutron;

� explain various rules for filling of electrons and write the distribution ofelectrons in different shells upto atomic number 20;

� define valency and correlate the electronic configuration of an atom withits valency;

� define atomic number and mass number of an atom;

� describe isotopes and isobars;

� define and compute average atomic mass and explain its fractional value.

5.1 CHARGED PARTICLES IN ATOM

You have read about Dalton’s atomic theory in lesson 3. The theory proposed inthe year 1803 considered the atom to be the smallest indivisible constituent of allmatter. The Dalton’s theory could explain the law of conservation of mass, law ofconstant composition and law of multiple proportions known at that time. However,towards the end of nineteenth century, certain experiments showed that an atom isneither the smallest nor indivisible particle of matter as stated by Dalton. It was shownto be made up of even smaller particles. These particles were called electrons,protons and neutrons. The electrons are negatively charged whereas the protons arepositively charged. The neutrons on the other hand are uncharged in nature. You willnow learn about the discovery of the charged subatomic particles.

5.1.1 Discovery of Electron

In 1885, Sir William Crookes carried out a series of experiments to study thebehaviour of metals heated in a vacuum using cathode ray tubes. A cathode ray tube

Fig. 5.1: A cathode ray tube; cathode rays are obtained on applying highvoltage across the electrodes in an evacuated glass tube

High voltage

95

Atomic Structure


Notes


consists of two metal electrodes in a partially evacuated glass tube. An evacuatedtube is the one from which most of the air has been removed. The negatively chargedelectrode is called cathode whereas the positively charged electrode is called anode.These electrodes are connected to a high voltage source. Such a cathode ray tubehas been shown in Fig. 5.1.

It was observed that when very high voltage was passed across the electrodes inevacuated tube, the cathode produced a stream of particles. These particles wereshown to travel from cathode to anode and were called cathode rays. In theabsence of external magnetic or electric field these rays travel in straight line. In 1897,an English physicist Sir J.J. Thomson showed that the rays were made up of a streamof negatively charged particles. This conclusion was drawn from the experimentalobservations when the experiment was done in the presence of an external electricfield. Following are the important properties of cathode rays:

� Cathode rays travel in straight line

� The particles constituting cathode rays carry mass and possess kinetic energy

� The particles constituting cathode rays have negligible mass but travel very fast

� Cathode ray particles carry negative charge and are attracted towards positivelycharged plate when an external electric field is applied (Fig. 5.2)

� The nature of cathode rays generated was independent of the nature of the gasfilled in the cathode ray tube as well as the nature of metal used for makingcathode and anode. In all the cases the charge to mass ratio (e/m) was foundto be the same.

+

B

A

Cathode Anode

Fluorescent screen

Fig. 5.2: The cathode rays are negatively charged; these travel in straight linefrom cathode to the anode (A), however in the presence of an external

electrical field these bend towards the positive plate (B)

These particles constituting the cathode rays were later called electrons. Since itwas observed that the nature of cathode rays was the same irrespective of the metalused for the cathode or the gas filled in the cathode ray tube. This led Thomson to

Atomic Structure


Notes


conclude that all atoms must contain electrons. This meant that the atom is notindivisible as was believed by Dalton and others. In other words, we can saythat the Dalton’s theory of atomic structure failed partially.

This conclusion raised a question, “If the atom was divisible, then what were itsconstituents?”. Today a number of smaller particles are found to constitute atoms.These particles constituting the atom are called subatomic particles. You have learntabove that electron is one of the constituents of the atom, let us study the next sectionto learn about another constituent particle present in an atom. As the atom is neutral,we expect the presence of positively charged particles in the atom so as to neutralisethe negative charge of the electrons.

5.1.2 Discovery of Proton

Much before the discovery of electron, Eugen Goldstein (in 1886) performed anexperiment using a perforated cathode (a cathode having holes in it) in the dischargetube filled with air at a very low pressure. When a high voltage was applied acrossthe electrodes in the discharge tube, a faint red glow was observed behind theperforated cathode. Fig. 5.3

Fig. 5.3 Goldstein’s cathode ray tube with perforated cathode

This glow was due to another kind of rays flowing in a direction opposite to thatof the cathode rays. These rays were called as anode rays or positive rays. Thesewere positively charged and were also called canal rays because they passedthrough the holes or the canals present in the perforated cathode. The followingobservations were made about anode rays (canal rays):

� Like cathode rays, the anode rays also travel in straight lines.

� The particles constituting anode rays carry mass and have kinetic energy.

� The particles constituting canal rays are much heavier than electrons and carrypositive charges

++

+++

+++

Red glow

Gas at LOWpressure

Positive raysfrom Anode

High voltage source

AnodePerforated cathode

+

97

Atomic Structure


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� The positive charge on the particles was whole number multiples of the amountof charge present on the electron.

� The nature and the type of the particles constituting the anode rays weredependent on the gas present in the discharge tube.

The origin of anode rays can be explained in terms of interaction of the cathode rayswith the gas present in the vacuum tube. It can be explained as given below:

The electrons emitted from the cathode collide with the neutral atoms of the gaspresent in the tube and remove one or more electrons present in them. This leavesbehind positive charged particles which travel towards the cathode. When thecathode ray tube contained hydrogen gas, the particles of the canal rays obtainedwere the lightest and their charge to mass ratio (e/m ratio) was the highest. Rutherfordshowed that these particles were identical to the hydrogen ion (hydrogen atom fromwhich one electron has been removed). These particles were named as protons andwere shown to be present in all matter. Thus, we see that the experiments by Thomsonand Goldstein had shown that an atom contains two types of particle which areoppositely charged and an atom is electrically neutral. What do you think is therelationship between the numbers of these particles in a given atom?

In addition to the two charged particles namely the electron and the proton, a neutralparticle called neutron was also discovered about which you would learn later in thislesson. Now, it is the time to check your understanding. For this, take a pause andsolve the following intext questions:

INTEXT QUESTIONS 5.1

1. Name two charged particles which constitute all matter.

2. Describe a cathode ray tube.

3. Name the negatively charged particles emitted from the cathode in the cathoderay tube?

4. Why do the canal rays obtained by using different gases have different e/m values?

In addition to the discovery of electrons and protons as the constituents ofatom, the phenomenon of radioactivity that is the spontaneous emission ofrays from atoms of certain elements also proved that the atom was divisible.

5.2 EARLIER MODELS OF ATOM

In section 5.1 you have learnt that the atom is divisible and contains three smallerparticles in it. The question that arises is, “In what way are the subatomic particles

Atomic Structure


Notes


arranged in the atom?”. On the basis of experimental observations, different modelshave been proposed for the structure of an atom. In this section, we will discusstwo such models namely Thomson model and Rutherford model.

5.2.1 Thomson Model

In lesson 3 you have learnt that all matter is made of atoms and all the atoms areelectrically neutral. Having discovered electron as a constituent of atom, Thomsonconcluded that there must be an equal amount of positive charge present in an atom.On this basis he proposed a model for the structure of atom. According to his model,atoms can be considered as a large sphere of uniform positive charge with a numberof small negatively charged electrons scattered throughout it, Fig. 5.4. This modelwas called as plum pudding model. The electrons represent the plums in the puddingmade of negative charge. This model is similar to a water-melon in which the pulprepresents the positive charge and the seeds denote the electrons. However, you maynote that a water melon has a large number of seeds whereas an atom may not haveas many electrons.

Electron

Sphere ofpositive charge

Fig. 5.4: Thomson’s plum-pudding model

5.2.2 Rutherford’s model

Ernest Rutherford and his co-workers were working in the area of radioactivity. Theywere studying the effect of alpha (α) particles on matter. The alpha particles are heliumnuclei, which can be obtained by the removal of two electrons from the helium atom.In 1910, Hans Geiger (Rutherford’s technician) and Ernest Marsden (Rutherford’sstudent) performed the famous α-ray scattering experiment. This led to the failureof Thomson’s model of atom. Let us learn about this experiment.

ααααα-Ray scattering experiment

In this experiment a stream of α particles from a radioactive source was directedon a thin (about 0.00004 cm thick) piece of gold foil. On the basis of Thomson’smodel it was expected that the alpha particles would just pass straight through the

99

Atomic Structure


Notes


gold foil and could be detected by a photographic plate placed behind the foil.However, the actual results of the experiment, Fig. 5.5, were quite surprising. It wasobserved that:

(i) Most of the α-particles passed straight through the gold foil.

(ii) Some of the α-particles were deflected by small angles.

(iii) A few particles were deflected by large angles.

(iv) About 1 in every 12000 particles experienced a rebound.

Beam of

-particles�

Thin gold foil

Most particlesare unreflected

Seattered of particles

Circularfluorescent

screen

Fig. 5.5: The experimental set-up and observations in the α- ray scatteringexperiment performed by Geiger and Marsden

The results of α-ray scattering experiment were explained by Rutherford in 1911 andanother model of the atom was proposed. According to Rutherford’s model, Fig. 5.6(a).

� An atom contains a dense and positively charged region located at its centre;it was called as nucleus,

� All the positive charge of an atom and most of its mass was contained in thenucleus,

� The rest of an atom must be empty space which contains the much smaller andnegatively charged electrons,

(a) (b)Fig 5.6: (a) Rutherford’s model of atom (b) Explanation of the results of scattering

experiment by Rutherford’s model.

-

+

Electron

Nucleus

--

-

-

-

-

- -

-

-

-

-

-+

Atomic Structure


Notes


On the basis of the proposed model, the experimental observations in the scatteringexperiment could be explained. This is illustrated in Fig. 5.6(b). The α particlespassing through the atom in the region of the electrons would pass straight withoutany deflection. Only those particles that come in close vicinity of the positively chargednucleus get deviated from their path. Very few α-particles, those that collide withthe nucleus, would face a rebound.

On the basis of his model, Rutherford was able to predict the size of the nucleus.He estimated that the radius of the nucleus was at least 1/10000 times smallerthan that of the radius of the atom. We can imagine the size of the nucleus withthe following analogy. If the size of the atom is that of a cricket stadium then thenucleus would have the size of a fly at the centre of the stadium.


1. Describe Thomson’s model of atom. What is it called?

2. What would have been observed in the α-ray scattering experiment if theThomson’s model was correct?

3. Who performed the α-ray scattering experiment and what were the observations?

4. Describe the model of atom proposed by Rutherford.

5.3 DRAWBACKS OF RUTHERFORD’S MODEL

According to Rutherford’s model the negatively charged electrons revolve in circularorbits around the positively charged nucleus. However, according to Maxwell’selectromagnetic theory (about which you may learn in higher classes), if a chargedparticle accelerates around another charged particle then it would continuously loseenergy in the form of radiation. The loss of energy would slow down the speed ofthe electron. Therefore, the electron isexpected to move in a spiral fashionaround the nucleus and eventually fall intoit as shown in Fig. 5.7. In other words,the atom will not be stable. However, weknow that the atom is stable and such acollapse does not occur. Thus,Rutherford's model is unable to explainthe stability of the atom. You know thatan atom may contain a number ofelectrons. The Rutherford’s model alsodoes not say anything about the way the electrons are distributed around the nucleus.Another drawback of the Rutherford’s model was its inability to explain the

Fig. 5.7: The electron in the Rutherford’smodel is expected to spiral into the nucleus

101

Atomic Structure


Notes


relationship between the atomic mass and atomic number (the number of protons).This problem was solved later by Chadwick by discovering neutron, the third particleconstituting the atom. You would learn about it in section 5.5.

The problem of the stability of the atom and the distribution of electrons in the atomwas solved by Neils Bohr by proposing yet another model of the atom. This isdiscussed in the next section.

5.4 BOHR’S MODEL OF ATOM

In 1913, Niels Bohr,a student of Rutherford proposed a model to account for theshortcomings of Rutherford’s model. Bohr’s model can be understood in terms oftwo postulates proposed by him. The postulates are:

Postulate 1: The electrons move in definite circular paths of fixed energy arounda central nucleus; just like our solar system in which different planets revolve aroundthe Sun in definite trajectory. Similar to the planets, only certain circular paths aroundthe nucleus are allowed for the electrons to move. These paths are called orbits,or energy levels. The electron moving in the orbit does not radiate. In other words,it does not lose energy; therefore, these orbits are called stationary orbits orstationary states. The bold concept of stationary state could answer the problemof stability of atom faced by Rutherford’s model.

32

K

18

L

8

M

2

N

+ Nucleus

(n=4) shellN

(n=3) shellM

(n=2) shellL

(n=1) shellK

Fig. 5.8: Illustration showing different orbits or the energy levels of fixedenergy in an atom according to Bohr’s model

It was later realised that the concept of circular orbit as proposed by Bohr was notadequate and it was modified to energy shells with definite energy. While a circularorbit is two dimensional, a shell is a three dimensional region. The shells of definiteenergy are represented by letters (K, L, M, N etc.) or by positive integers (1, 2,3, …. etc.) Fig. 5.8. The energies of the shells increase with the number n; n = 1,

Atomic Structure


Notes


level is of the lowest energy. Further, the maximum number of electrons that can beaccommodated in each shell is given by 2n2, where n is the number of the level. Thus,the first shell (n=1) can have a maximum of two electrons whereas the second shellcan have 8 electrons and so on. Each shell is further divided into various sublevelscalled subshells about which you would study in your higher classes.

Postulate 2: The electron can change its shells or energy level by absorbing orreleasing energy. An electron at a lower state of energy Ei can go to a final higherstate of energy Ef by absorbing a single photon of energy given by:

E = hν = Ef – Ei

Similarly, when electron changes its shell from a higher initial level of energy Ei toa lower final level of energy Ef, a single photon of energy hν is released (Fig. 5.9).

n = 3

n = 2

n = 1

A photon is emitted

with energy E=h�

Increasing energyof orbits

Fig. 5.9: The electrons in an atom can change their energy level by absorbingsuitable amounts of energy or by emitting energy.


1. Give any two drawbacks of Rutherford’s model of atom.

2. State the postulates of Bohr’s model.

3. How does Bohr model of an atom explain the stability of the atom?

Thus, the Bohr’s model of atom removes two of the limitations of Rutherford’s model.These are related to the stability of atom and the distribution of electrons aroundthe nucleus. You would recall that the third limitation of Rutherford’s model was its

103

Atomic Structure


Notes


inability to explain the relationship between the atomic mass, and the atomic number(the number of protons) of an atom. Let us learn how this problem was solved withthe discovery of neutron.

5.5 DISCOVERY OF NEUTRON

You would recall that when we discussed about the failure of Rutherford’s modelwe mentioned that it was unable to explain the relationship between the atomic massand the atomic number (the number of protons). According to the Rutherford’s model,the mass of helium atom (containing 2 protons) should be double that of a hydrogenatom (with only one proton). [Ignoring the mass of electron as it is very light].However, the actual ratio of the masses of helium atom to hydrogen atom is 4:1.It was suggested that there must be one more type of subatomic particle presentin the nucleus which may be neutral but have mass.

Such a particle was discovered by James Chadwick in 1932. This was found to beelectrically neutral and was named neutron. Neutrons are present in the nucleus ofall atoms, except hydrogen. A neutron is represented as ‘n’ and is found to havea mass slightly higher than that of a proton. Thus, if the helium atom contained 2protons and 2 neutrons in the nucleus, the mass ratio of helium to hydrogen (4:1)could be explained. The characteristics of the three fundamental particles constitutingthe atom are given in Table 5.1.

Table 5:1 Characteristics of the fundamental subatomic particles

Particle Symbol Mass (in kg) Actual Charge Relative charge(in Coulombs)

Electron e 9.109 389 × 10–31 1.602 177 × 10–19 –1

Proton p 1.672 623 × 10–27 1.602 177 × 10–19 1

Neutron n 1.674 928 × 10–27 0 0


1. What is a neutron and where is it located in the atom?

2. How many neutrons are present in the α-particle?

3. How will you distinguish between an electron and a proton?

5.6 ATOMIC NUMBER AND MASS NUMBER

You have learnt that the nucleus of atom contains positively charged particles calledprotons and neutral particles called neutrons. The number of protons in an atomis called the atomic number and is denoted by the symbol ‘Z’. All atoms of an

Atomic Structure


Notes


element have the same atomic number. The electrons occupy the space outside thenucleus. In order to account for the electrically neutral nature of the atom, the numberof protons in the nucleus is exactly equal to the number of electrons. Thus,

Atomic number = number of protons = number of electrons

You would recall that according to Dalton’s theory, the atoms of different elementsare different from each other. We can now say that this difference is due to differencein the numbers of protons present in the nucleus of the element. In other words,different elements differ in terms of their atomic number. For example, the atomsof hydrogen and helium are different because hydrogen has one proton in its nucleuswhereas the nucleus of helium atom contains two protons. Their atomic numbers are1 and 2, respectively. You have learnt in the Rutherford’s model that the mass ofthe atom is concentrated in its nucleus. This is due to the presence of two heavyparticles namely protons and neutrons in the nucleus. These particles are callednucleons. The number of nucleons in the nucleus of an atom is called its massnumber. It is denoted by ‘A’ and is equal to the total number of protons and neutronspresent in the nucleus of an element. Thus,

Mass number (A) = number of protons(Z) + number of neutrons(n)

Atomic number and mass number are represented on the symbol of an element. Anelement, X with an atomic number, Z and the mass number, A is denoted as follows:

AX

Z

For example, 12

C6

means that the carbon has an atomic number of 6 and the mass

number of 12. This can be used to compute the number of different fundamentalparticles in the atom. Let us calculate it for carbon.

As the atomic number is 6 this means:

Number of protons = number of electrons = 6

As Mass number = number of protons + number of neutrons

⇒ 12 = 6 + number of neutrons

⇒ number of neutrons = 12 – 6 = 6

Thus, an atom of 12

C6

has 6 protons, 6 electrons and 6 neutrons.


1. A sodium atom has an atomic number of 11 and a mass number of 23. Calculatethe number of protons, electrons and neutrons in a sodium atom.

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Atomic Structure


Notes


2. What is the mass number of an atom which has 7 protons and 8 neutrons?

3. Calculate the number of electrons, protons and neutrons in 40

Ar18

and 49

K19

.

5.7 ELECTRONIC CONFIGURATION: DISTRIBUTIONOF ELECTRONS IN DIFFERENT ORBITS

As discussed in section 5.4, the electrons move in definite paths called orbits or shellsaround a central nucleus. These orbits or shells have different energies and canaccommodate different number of electrons in them. The question arises that howare the electrons distributed amongst these shells? The answer to this question wasprovided by Bohr and Bury. According to their scheme, the electron distribution isgoverned by the following rules:

I. These orbits or shells in an atom are represented by the letters K, L, M, N,…or the positive integral numbers, n = 1,2,3,4,….

II. The orbits are arranged in the order of increasing energy. The energy of M shellis more than that of the L shell which in turn is more than that of the K shell.

III. The maximum number of electrons present in a shell is given by the formula 2n2,where ‘n’ is the number of the orbit or the shell. Thus, the maximum numberof electrons that can be accommodated in different shells are as follows:

Maximum number of electrons in K shell (or n = 1 level) = 2n2 = 2 × (1)2

= 2

Maximum number of electrons in L shell (or n = 2 level) = 2n2 = 2 × (2)2

= 8

Maximum number of electrons in M shell (or n = 3 level) = 2n2 = 2 × (3)2

= 18 and so on. See table 5.2

Table 5.2: Electron accommodation capacity of different shells

Value of n Shell name Maximum capacity

1 K-Shell 2

2 L- Shell 8

3 M- Shell 18

4 N- Shell 32

IV. The shells are occupied in the increasing order of their energies.

V. Electrons are not accommodated in a given shell, unless the inner shells arecompletely filled.

The arrangement of electrons in the various shells or orbits of an atom of the elementis known as electronic configuration. Keeping these points in mind, let us now studythe filling of electrons in various shells of atoms of different elements.

Atomic Structure


Notes


� Hydrogen (H) atom has only one electron. It would occupy the first shell andelectronic configuration of hydrogen can be represented as 1.

� The next element helium (He) has two electrons in its atom. Since the first shellcan accommodate two electrons; hence, this second electron will also be placedin the first shell. The electronic configuration of helium is written as 2.

� The third element, Lithium (Li) has three electrons. Now the two electrons occupythe first shell whereas the third electron goes to the next shell of higher energylevel, i.e. second shell. Thus, the electronic configuration of Li is 2, 1.

Similarly, the electronic configurations of other elements can be written. The structuresof the atoms of elements with atomic number 1 to 18 are given in Fig. 5.10.

H

Li

Na

Be

Mg

B

Al

C

Si

N

P

O

S

F

Cl

Ne

Ar

He

Fig. 5.10: The structures, according to Bohr’s model of atoms, of elementswith atomic number 1 to 18.

5.7.1 Concept of Valence or Valency

We have just discussed the electronic configuration of first 18 elements. You can seefrom the Fig. 5.10 that different elements have different number of electrons in theoutermost or the valence shell. These electrons in the outermost shell are known asvalence electrons. The number of valence electrons determines the combiningcapacity of an atom in an element. Valence is the number of chemical bonds thatan atom can form with univalent atoms. Since hydrogen is a univalent atom, thevalence of an element can be taken by the number of atoms of hydrogen with whichone atom of the element can combine. For example, in H2O, NH3, and CH4 thevalencies of oxygen, nitrogen and carbon are 2, 3 and 4 respectively.

The elements having a completely filled outermost shell in their atoms show little orno chemical activity. In other words, their combining capacity or valency is zero. Theelements with completely filled valence shells are said to have stable electronic

107

Atomic Structure


Notes


configuration. The main group elements can have a maximum of eight electrons intheir valence shell. This is called octet rule; you will learn more about it in lesson 7.You will learn that the combining capacity or the tendency of an atom to react withother atoms to form molecules depends on the ease with which it can achieve octetin its outermost shell. The valencies of the elements can be calculated from theelectronic configuration by applying the octet rule. It can be seen as follows:

� If the number of valence electrons is four or less then the valency is equal tothe number of the valence electrons.

� In cases when the number of valence electrons is more than four then generallythe valency is equal to 8 minus the number of valence electrons.

Thus,

Valency = Number of valence electrons (for 4 or lesser valence electrons)

Valency = 8 - Number of valence electrons (for more than 4 valence electrons)

The composition and electronic configuration of the elements having the atomicnumbers from 1 to 18, along with their valencies is given in Table 5.3.

Table 5.3: The composition, electron distribution and common valencyof the elements with atomic number from 1 to 18

Hydrogen

Helium

Lithium

Beryllium

Boron

Carbon

Nitrogen

Oxygen

Neon

Sodium

Fluorine

Magnesium

Aluminium

Silicon

Phosphorus

Sulphur

Chlorine

Argon

H

He

Li

Be

B

C

N

O

F

Ne

Na

Mg

Al

Si

P

S

Cl

Ar

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

–

2

4

5

6

6

7

8

10

10

12

12

14

14

16

16

18

22

1

2

2

2

2

2

2

2

2

2

2

2

2

2

2

2

2

2

–

–

1

2

3

4

5

6

7

8

8

8

8

8

8

8

8

8 8

7

6

5

4

3

2

1

–

–

–

–

–

–

––

–

–

–

–

–

–

–

–

––

–

–

––

–

–

–

–

–

–

1

1

1

1

1

2

2

2

2

3, 5

3

3

3

4

0

0

0

4

Name ofElement

Symbol

K

AtomicNumber

Numberof

Protons

Numberof

Neutrons

Numberof

Electrons

Distribution ofElectrons

L M N

Valency

Atomic Structure


Notes


In next lesson, you will study about the importance of electronic configurations inunderstanding the periodic arrangement of elements. These electronic configurationsare also helpful in studying the nature of bonding between various elements whichwill be dealt with in lesson 7.


1. How many shells are occupied in the nitrogen (atomic number =7) atom?

2. Name the element which has completely filled first shell.

3. Write the electronic configuration of an element having atomic number equal to 11.

WHAT HAVE YOU LEARNT

� According to Dalton’s atomic theory, the atom is considered to be the smallestindivisible constituent of all matter. This theory could explain the law ofconservation of mass, law of constant composition and law of multiple proportions.However, certain experiments towards the end of nineteenth century showedthat the atom is neither the smallest nor indivisible particle of matter. It was shownto be made up of even smaller particles called electrons, protons and neutrons.

� Sir J.J.Thomson discovered that when very high voltage was passed across theelectrodes in the cathode ray tube, the cathode produced rays that travel fromcathode to anode and were called cathode rays. It showed that the rays weremade up of a stream of negatively charged particles called electrons. Thediscovery of electrons meant that the atom is not indivisible as was believedby Dalton and others.

� Eugen Goldstein discovered anode rays by using a perforated cathode (acathode having holes in it) in the discharge tube filled with air at a very lowpressure. The discovery of anode rays established the presence of positivelycharged proton in the atom.

� According to Thomson’s plum-pudding model, atoms can be considered as alarge sphere of uniform positive charge with a number of small negatively chargedelectrons scattered throughout it.

� The α-ray scattering experiment performed by Geiger and Marsden led to thefailure of Thomson’s model of atom. In this experiment, a stream of α-particlesfrom a radioactive source was directed on a thin piece of gold foil. Most of theα-particles passed straight through the gold foil, some α-particles were deflectedby small angles, a few particles by large angles and very few experienced arebound.

109

Atomic Structure


Notes


� The results of α-ray scattering experiment were explained in terms of Rutherford’smodel. According to which the atom contains a dense and positively chargedregion called nucleus at its centre and the negatively charged electrons movearound it. All the positive charge and most of the mass of atom is contained inthe nucleus.

� The Rutherford’s model however failed as it could not explain the stability ofthe atom, the distribution of electrons and the relationship between the atomicmass and atomic number (the number of protons).

� The problem of the stability of the atom and the distribution of electrons in theatom was solved by Neils Bohr in terms of Bohr’s model of the atom. Bohr’smodel can be understood in terms of two postulates, the first being, ‘ Theelectrons move in definite circular paths of fixed energy around a centralnucleus’ and the second, ‘The electron can change its orbit or energy levelby absorbing or releasing energy.’

� In 1932, James Chadwick discovered an electrically neutral particle in atom andnamed it as neutron.

� The number of protons in an atom is called the atomic number and is denotedas ‘Z’. On the other hand the number of nucleons( protons plus neutrons) inthe nucleus of an atom is called its mass number and is denoted as ‘A’

� The electrons are distributed in different shells in the order of increasing energy.The distribution is called electronic configuration. The maximum number ofelectrons present in a shell is given by the formula 2n2, where ‘n’ is the numberof the orbit or the shell.

� The valence is the number of chemical bonds that an atom can form with univalentatoms. If the number of valence electrons is four or less, then the valency is equalto the number of the valence electrons. On the other hand, if the number ofvalence electrons is more than four, then generally the valency is equal to 8 minusthe number of valence electrons.

TERMINAL EXERCISE

1. How did J.J.Thomson discover the electron? Explain his “plum pudding” modelof the atom.

2. What made Thomson conclude that all atoms must contain electrons?

3. Identify the following subatomic particles:

(a) The number of these in the nucleus is equal to the atomic number

(b) The particle that is not found in the nucleus

(c) The particle that has no electrical charge

(d) The particle that has a much lower mass than the others subatomic particles

Atomic Structure


Notes


4. Which of the following are usually found in the nucleus of an atom?

(a) Protons and neutrons only

(b) Protons, neutrons and electrons

(c) Neutrons only

(d) Electrons and neutrons only

5. Describe Ernest Rutherford’s experiment with alpha particles and gold foil. Howdid this lead to the discovery of the nucleus?

6. What does the atomic number tell us about an atom?

7. What is the relationship between the numbers of electrons and protons in anatom?

8. How did Neils Bohr revise Rutherford’s atomic model?

9. What is understood by a stationary state?

10. What is a shell? How many electrons can be accomodate in L-shell?

11. State the rules for writing the electronic configuration of elements.

ANSWERS TO INTEXT QUESTIONS

5.1

1. Electrons and protons

2. A cathode ray tube consists of two metal electrodes in a partially evacuated glasstube. The negatively charged electrode is called cathode while the positivelycharged electrode is called anode. These electrodes are connected to a highvoltage source.

3. Electron

4. When the electrons emitted from the cathode collide with the neutral atoms ofthe gas present in the tube, these remove one or more electrons present in them.This leaves behind positive charged particles which travel towards the cathode.As the atoms of different gases have different number of protons present in them,these give positively charged ions with different e/m values.

5.2

1. According to Thomson’s model, atoms can be considered as a large sphere ofuniform positive charge with a number of small negatively charged electronsscattered throughout it. This model was called as plum pudding model.

111

Atomic Structure


Notes


2. If the Thomson’s model was correct, then most of the α-particles in the α-rayscattering experiment would have passed straight through the atom

3. The α-ray scattering experiment was performed by Geiger and Marsden. Whena stream of α-particles from a radioactive source was directed on a thin pieceof gold foil, most of the α-particles passed straight through the gold foil, someα-particles were deflected by small angles, a few particles by large angles andvery few experienced a rebound.

4. According to Rutherford’s model, the atom contains a dense and positivelycharged region called nucleus at its centre and the negatively charged electronsmove around it. All the positive charge and most of the mass of atom is containedin the nucleus.

5.3

1. The Rutherford’s model could not explain the stability of the atom, the distributionof electrons and the relationship between the atomic mass and atomic number(the number of protons).

2. The two postulates of Bohr’s model are :

I. The electrons move in definite circular paths of fixed energy around acentral nucleus.

II. The electron can change its orbit or energy level by absorbing or releasingenergy.

3. The Bohr’s model explains the stability of atom by proposing that the electrondoes not lose energy when present in a given energy level.

5.4

1. It is a neutral subatomic particle present in the nucleus of the atom.

2. An α-particle contains two neutrons.

3. The electron and proton can be distinguished in terms of their charge and mass.While the electron is negatively charged, the proton is positively charged.Secondly, the proton is much heavier than the electron; it is about 1840 timesheavier.

5.5

1. No of protons = 11

No. of electrons = 11

No. of neutrons = 12

Atomic Structure


Notes


2. Mass number = number of protons + number of neutrons

Therefore, mass number = 7 + 8 = 15

3.40

Ar18

: Number of protons = atomic number = 18

Number of electrons = number of protons = 18

Number of neutrons = mass number – number of protons = 40 – 18 = 22

40K

19 Number of protons = atomic number = 19

Number of electrons = number of protons = 19

Number of neutrons = mass number – number of protons = 40 – 19 = 21

5.6

1. The electronic configuration of nitrogen is 2, 5. Thus, two shells are occupied.The first shell (capacity = 2) is completely filled while the second shell (capacity= 8) is partially filled.

2. Helium

3. The electronic configuration of an element having atomic number 11 is 2, 8, 1.

Education

ATOMIC STRUCTURE